Copyright © Houghton Mifflin Company. All rights reserved.17a–1

Copyright © Houghton Mifflin Company. All rights reserved.17a–2 Electrochemistry The study of the interchange of chemical and electrical energy.

Copyright © Houghton Mifflin Company. All rights reserved.17a–3 Half-Reactions The overall reaction is split into two half- reactions, one involving oxidation and one reduction. 8H + + MnO 4  + 5Fe 2+  Mn Fe H 2 O Reduction: 8H + + MnO 4  + 5e   Mn H 2 O Oxidation: 5Fe 2+  5Fe e 

Copyright © Houghton Mifflin Company. All rights reserved.17a–4 Galvanic Cell A device in which chemical energy is changed to electrical energy.

Copyright © Houghton Mifflin Company. All rights reserved.17a–5 Anode and Cathode OXIDATION occurs at the ANODE. REDUCTION occurs at the CATHODE.

Copyright © Houghton Mifflin Company. All rights reserved.17a–6 Figure 17.1: Schematic of a method to separate the oxidizing and reducing agents of a redox reaction. (The solutions also contain counterions to balance the charge.)

Copyright © Houghton Mifflin Company. All rights reserved.17a–7 Figure 17.2: Galvanic cells can contain a salt bridge as in (a) or a porous-disk connection as in (b).

Copyright © Houghton Mifflin Company. All rights reserved.17a–8 Figure 17.3: An electrochemical process involves electron transfer at the interface between the electrode and the solution. (a) The species in the solution acting as the reducing agent supplies electrons to the anode. (b) The species in the solution acting as the oxidizing agent receives electrons from the cathode.

Copyright © Houghton Mifflin Company. All rights reserved.17a–9 Figure 17.4: Digital voltmeters draw only a negligible current and are convenient to use.

Copyright © Houghton Mifflin Company. All rights reserved.17a–10 Figure 17.5: (a) A galvanic cell involving the reactions Zn Zn2+ + 2e- (at the anode) and 2H+ + 2e- H 2 (at the cathode) has a potential of 0.76 V. (b) The standard hydrogen electrode where H 2 (g) at 1 atm is passed over a platinum electrode in contact with 1 M H+ ions. This electrode process (assuming ideal behavior) is arbitrarily assigned a value of exactly zero volts.

Copyright © Houghton Mifflin Company. All rights reserved.17a–11 Cell Potential Cell Potential or Electromotive Force (emf): The “pull” or driving force on the electrons.

Copyright © Houghton Mifflin Company. All rights reserved.17a–12 An electrochemical cell with a measured potential of 1.10 V.

Figure 17.6: A galvanic cell involving the half-reactions Zn Zn e- (anode) and Cu e- Cu (cathode), with  ºcell = 1.10 V.

Copyright © Houghton Mifflin Company. All rights reserved.17a–14

Copyright © Houghton Mifflin Company. All rights reserved.17a–15 Figure 17.7: The schematic of a galvanic cell based on the half-reactions

Copyright © Houghton Mifflin Company. All rights reserved.17a–16 Figure 17.8: Schematic diagram for the galvanic cell based on the half-reactions

Copyright © Houghton Mifflin Company. All rights reserved.17a–17 Figure 17.9: A concentration cell that contains a silver electrode and aqueous silver nitrate in both compartments.

Copyright © Houghton Mifflin Company. All rights reserved.17a–18 Figure 17.10: A concentration cell containing iron electrodes and different concentrations of Fe 2+ ion in the two compartments.

Copyright © Houghton Mifflin Company. All rights reserved.17a–19 Figure 17.19: (a) A standard galvanic cell based on the spontaneous reaction (b) A standard electrolytic cell. A power source forces the opposite reaction

Copyright © Houghton Mifflin Company. All rights reserved.17a–20 Standard Reduction Potentials The E  values corresponding to reduction half-reactions with all solutes at 1M and all gases at 1 atm. Cu e   Cu E  = 0.34 V vs. SHE SO 4 2  + 4H + + 2e   H 2 SO 3 + H 2 O – E  = 0.20 V vs. SHE

Copyright © Houghton Mifflin Company. All rights reserved.17a–21 Free Energy and Cell Potential  G  =  nFE  n = number of moles of electrons F = Faraday = 96,485 coulombs per mole of electrons

Copyright © Houghton Mifflin Company. All rights reserved.17a–22 The Nernst Equation We can calculate the potential of a cell in which some or all of the components are not in their standard states.

Copyright © Houghton Mifflin Company. All rights reserved.17a–23 Calculation of Equilibrium Constants for Redox Reactions At equilibrium, E cell = 0 and Q = K.

Copyright © Houghton Mifflin Company. All rights reserved.17a–24 Fuel Cells...galvanic cells for which the reactants are continuously supplied. 2H 2 (g) + O 2 (g)  2H 2 O(l) anode: 2H 2 + 4OH   4H 2 O + 4e  cathode: 4e  + O 2 + 2H 2 O  4OH 

Copyright © Houghton Mifflin Company. All rights reserved.17a–25 Electrolysis...forcing a current through a cell to produce a chemical change for which the cell potential is negative.

Copyright © Houghton Mifflin Company. All rights reserved.17a–26